Molecule Gallery - Aromatic Rings Molecule Gallery Aromatic Rings Aromatic rings (also known as aromatic compounds or arenes) are hydrocarbons which contain benzene, or some other related ring structure . Benzene, C6H6, is often drawn as a ring of six carbon atoms, with alternating double bonds and single bonds: This simple picture has some complications, however. (Everything in organic chemistry has complications!) Carbon-carbon single bonds are longer than carbon-carbon double bonds, so if there were "real" single and double bonds in the molecule, the shape of the benzene molecule would be a distorted hexagon: What is actually found is that all of the bond lengths in the benzene rings are 1.397 angstroms, which is roughly intermediate between the typical lengths of single bonds (~1.5 angstroms) and double bonds (~1.3 angstroms). All of the carbon atoms in the benzene rings are sp2-hybridized: the overlap of the sp2 orbitals around the ring produces a framework of six sigma bonds, while the unhybridized p-orbitals which are perpendicular to this plane overlap in a side-to-side fashion to form three pi-bonds.

These pi-bonds are delocalized around the ring, leading to an unusual stability for the benzene ring compared to other alkenes. Instead of three "real" single bonds and three "real" double bonds in the ring, benzene may be thought of as consisting of "six one-and-a-half" bonds, represented by the resonance structures shown below: All of the positions on the benzene ring are thus equivalent; benzene is sometimes represented as a hexagon with a circle inside it to emphasize this equivalence: Since all of the atoms in the ring are sp2-hybridized, they are all trigonal planar, with bond angles of 120ï¿½, and the benzene ring is a flat molecule, shaped like a hexagon. Aromatic hydrocarbons are nonpolar, and are insoluble in water. However, when other atoms are substituted on the benzene ring, they may be very water-soluble. For instance, phenol , which has an —OH group attached to the benzene ring, is very water-soluble.

Because of the stability imparted by the delocalized pi-electrons in the ring, aromatic molecules do not undergo many of the reactions which are typical for alkenes, but there are a number of interesting and useful substitution reactions that they undergo, in which hydrogen atoms on the ring are replaced with other functional groups. Molecules which do not contain benzene rings are referred to aliphatic ("fat-like") compounds. Aromatic molecules with simple alkyl groups as substituents are named as derivatives of benzene. For instance, a benzene with an ethyl group attached to one of the carbons in the ring is simply called "ethylbenzene." A number of aromatic molecules are known by common names; for instance, benzene with a —CH3 group attached is called " toluene "; benzene with an —NH2 group attached is called " aniline "; a benzene with a —CO2H group is called "benzoic acid," etc.

For more complicated substituents, the benzene ring is named as a substituent, in which case it is called "phenyl." For instance, an eight-carbon chain with a benzene ring on the third carbon is called "3-phenyloctane." When there are two substituents on the benzene ring, numbers can be used to identify the position of the substituents, but an older system of nomenclature is often used instead, in which the prefixes ortho-, meta-, and para- (often abbreviated as o-, m-, and p-) are used to indicate the relative placement of the substituents: For three or more substituents, numbers must be used to indicate the placement of substituents. Annulenes 4n+2 rule Molecules whose Lewis structures can be drawn as a ring with alternating single and double bonds are known as annulenes. Annulenes are named by prefixing the number of carbon atoms in the rings in brackets before the word annulene: hence, benzene can be named as [6] annulene.

Hï¿½ckel's Rule (Erich Hï¿½ckel, 1931) states that annulenes with have (4n + 2) p electrons (where n is an integer from 0, 1, 2, 3, ...) are resonance-delocalized in the same way that benzene is, and should be aromatic. Thus, molecules with 2, 6, 10, 14, 18, etc., p electrons will be aromatic, while annulenes with 4, 8, 12, 16, 20, etc. p electrons will not be aromatic, and their chemistry will be more similar to that of alkenes than to benzene. Chlorophyll 3D Chlorophyll (Greek, chloros "green" and phyllon "leaf") is a porphyrin derivative found in green plants and cyanobacteria which allows these organisms to perform photosynthesis. Shown below are the structures of chlorophyll a and chlorophyll b; the magnesium-free forms of these molecules are called pheophytin a and pheophytin b.

(The three-dimensional structure shown on the right is that of pheophytin a.) A double bond in one of the pyrrole rings in the porphyrin macrocycle is reduced to a single bond (in the structure below, this is the carbon-carbon bond to which the long side-chain, sometimes called the phytyl group, is attached); these "dihydroporphyrins" are known as chlorins. Chlorophyll absorbs visible light in the violet (~420 nm) and red (~680 nm) regions of the electromagnetic spectrum; this renders the light that is reflected off of this pigment to be green.

The energy from the sunlight that is absorbed by chlorophyll in the chloroplasts of green plants is used to drive the synthesis of carbohydrates such as glucose (C6H12O6) from carbon dioxide and water: (This equation summarizes a lot of complex chemistry, which it is beyond the scope of this page to try to summarize.) In the fall, the green color of many plants disappears as the chlorophyll starts to break down; this allows other pigments, such as carotenes, in the leaves to "show their colors," producing the vivid yellows and oranges associated with fall colors. Chlorophyll a References P. W. Atkins, Molecules, 2nd ed. Cambridge: Cambridge University Press, 2003, p. 46-48, 143, 144-151, 168-169. Paula Yurkanis Bruice, Organic Chemistry, 4th ed. Upper Saddle River: Prentice Hall, 2004, p. 594-599. Marye Anne Fox and James K. Whitesell, Organic Chemistry, 3rd ed.

Sudbury: Jones and Bartlett Publishers, p. 64-74. Maitland Jones, Jr., Organic Chemistry. New York: W. W. Norton & Company, 1997, p. 568-602. Richard J. Lewis, Sr., Hawley's Condensed Chemical Dictionary, 13th ed. New York: Van Nostrand Reinhold, 1997. Robert Thornton Morrison and Robert Neilson Boyd, Organic Chemistry, 6th ed. Englewood Cliffs: Prentice Hall, 1992, p. 493-513. Royston M. Roberts, Serendipity: Accidental Discoveries in Science. New York: John Wiley & Sons, Inc., 1989, p. 75-78. D. W. A. Sharp, The Penguin Dictionary of Chemistry, 2nd ed. London: Penguin Books, 1990. Graham Solomons and Craig Fryhle, Organic Chemistry, 7th ed. New York: John Wiley & Sons, Inc., 2000, p.

53-54, 627-643. L. G. Wade, Jr., Organic Chemistry, 5th ed. Upper Saddle River: Prentice Hall, 2003, p. 69, 679-706. Martha Windholz (ed.), The Merck Index, 10th ed. Rahway: Merck & Co., Inc., 1983. The Benzene Molecule 3D/Jmol The Benzene Molecule The chemical compound benzene (C6H6) is a colorless, flammable, aromatic hydrocarbon, that is a known carcinogen. It boils at 80.1°C and melts at 5.5°C. Benzene has a heat of vaporization of 44.3 kJ/mol and a heat of fusion of 9.84 kJ/mol. Produced by hydrogen reduction of some allotropes of carbon , or from petroleum , it is used in the creation of drugs , plastics, gasoline , synthetic rubber, napalm and dyes.

Benzene Molecule using Jsmol see -- For 3D Structures Structure of the Benzene Molecule The formula of benzene (C6H6), caused a mystery for some time after its discovery, as no proposed structure could take account of all the bonds (carbon usually forms four single bonds and hydrogen one). The chemist Kekulé was the first to deduce the ring structure of benzene; after years of studying carbon bonding, benzene and related molecules, the solution to the benzene structure came to him in a dream of a snake eating its own tail. Upon waking was inspired to deduce the ring structure of benzene. While his claims were well publicized and accepted, by the early 1920s Kekulé's biographer came to the conclusion that Kekulé's understanding of the tetravalent nature carbon bonding depended on the previous research of Archibald Scott Couper (1831-1892); further, the German Chemist Josef Loschmidt (1821-1895) had earlier posited a cyclic structure for benzene as early as 1862, although he had not actually proved this structure to be correct.

Benzene presents a problem, as to account for all the bonds in the molecule, there must be alternating double carbon bonds: Benzene Molecular Structure However, all of the carbon-carbon bonds in the benzene molecule are of the same length, and it is known that a single bond is longer than a double bond. In addition, the bond length (the distance between the two bonded atoms) in benzene is greater than a double bond, but shorter than a single bond. There seems in effect to be a bond and a half between each carbon. This is explained by electron delocalization. In order to picture this, we must consider the position of electrons in the bonds of benzene. One representation is that the structure of the benzene molecule exists as a superposition of the forms below, rather than either form individually. This type of structure is called a resonance hybrid of the benzene molecule. In reality, neither form really exists. Delocalisation must be explained using a higher level of theory than single and double bonds. The single bonds are formed with electrons in line between the carbon atoms - this is called Ã (sigma) symmetry. Double bonds consist of a sigma bond and another, À (pi) bond.

This second bond has electrons orbiting in paths above and below the plane of the ring at each bonded carbon atom. The À-bonds are formed from atomic p-orbitals above and below the plane of ring. The following diagram shows the positions of these p-orbitals in the benzene molecule: Being out of the plane of the atoms, these orbitals can interact with each other freely, and become delocalised. This means that instead of being tied to one atom of carbon, each electron is shared by all six in the ring. Thus there are not enough to form double bonds on all the carbon atoms, but the "extra" electrons do strengthen all of the bonds on the ring equally. The resulting molecular orbital has À symmetry. This delocalisation of electrons is known as aromaticity, and gives benzene great stability. This is the fundamental property of aromatic chemicals which differentiates them from non aromatics. To reflect the delocalised nature of the bonding, the benzene molecule may be depicted as a circle inside a hexagon in chemical structure diagrams. As with most diagrams of molecular structures, the Hydrogen atoms are frequently omitted.

Substituents Many important chemicals are essentially benzene, with one or more of the hydrogen atoms replaced with some other functionality, eg: Hydrocarbon Substituents Heterocycles - Carbons in the ring are replaced with something else. pyridine- is the most commonly encountered heterocycle, due to its aromaticity furan- is not stable, due to its electronic configuration; it is not aromatic. The derivative tetrahydrofuran (THF) is very widely used as a solvent and reagent. Production Benzene may result whenever carbon-rich materials undergo incomplete combustion. It is produced naturally in volcanoes and forest fires, and is also a component of cigarette smoke. Industrially, it is produced from either coal or petroleum. The steel industry 'cokes' coal. This coke is then 'cracked' to yield Benzene(63%), Toluene(14%) & Xylene(7%). Alternatively an Olefin plant will produce Benzene as a by-product of cracking naptha or gas oil. It is one of the components of the coal tars given off when coal in converted to coke. Up until World War II, this source of benzene was sufficient to meet world demand for the chemical.

However, in the 1950's, increasing demand for benzene, especially for the growing plastics necessitated the production of benzene from petroleum. Today, most benzene comes from the petrochemical industry, with only a small fraction being produced from coal. Uses Prior to the 1920's, benzene was frequently used as an industrial solvent, especially for degreasing metal. As its toxicity became obvious, other solvents replaced benzene in application that directly exposed the user to benzene. As a gasoline additive, benzene increases the octane rating and reduces knocking. As a result, gasoline once often contained several percent benzene, although in the 1950s tetraethyl lead replaced it as the most widely used antiknock additive. However, with the global phaseout of leaded gasoline, benzene has made a comeback as a gasoline additive in some nations. In the United States, concern over its negative health effects and the possibility of benzene entering the groundwater have led to stringent regulation of gasoline's benzene content. Many European nations, on the other hand, either do not regulate benzene in gasoline or allow significant quantities of it, and European gasoline formulations often contain 5% or more benzene.

By far the largest use of benzene is an intermediate to make other chemicals. The most widely produced derivatives of benzene are styrene, which is used to make polymers and plastics phenol for resins and adhesives (via cumene), and cyclohexane, which is used in Nylon manufacture. Smaller amounts of benzene are used to make some types of rubbers, lubricants, dyes, detergents, drugs, explosives and pesticides. Health Hazards Breathing very high levels of benzene can result in death, while high levels can cause drowsiness, dizziness, rapid heart rate, headaches, tremors, confusion, and unconsciousness. Eating or drinking foods containing high levels of benzene can cause vomiting, irritation of the stomach, dizziness, sleepiness, convulsions, rapid heart rate, and death. The major effect of benzene from chronic (long term) exposure is to the blood. Benzene causes harmful effects on the bone marrow and can cause a decrease in red blood cells leading to anemia. It can also cause excessive bleeding and can affect the immune system, increasing the chance of infection.

Some women who breathed high levels of benzene for many months had irregular menstrual periods and a decrease in the size of their ovaries. It is not known whether benzene exposure affects the developing fetus in pregnant women or fertility in men. Animal studies have shown low birth weights, delayed bone formation, and bone marrow damage when pregnant animals breathed benzene. Structure of Benzene Electrophilic Aromatic Substitution -- Mechanism Today we'll find that resonance is very important in understanding both the structure and the reactions of aromatic compounds. First, let's take a look at the structural representations which distinguish aromatic compounds from those that aren't aromatic. The most commonly encountered aromatic compound is benzene. The usual structural representation for benzene is a six carbon ring (represented by a hexagon) which includes three double bonds. Each of the carbons represented by a corner is also bonded to one other atom. In benzene itself, these atoms are hydrogens. The double bonds are separated by single bonds so we recognize the arrangement as involving conjugated double bonds. An alternative symbol uses a circle inside the hexagon to represent the six pi electrons. Each of these symbols has good and bad features.

We'll use the three double bond symbol simply because it is also routinely used in the text. Keep in mind that if the hexagon contains neither the three double bonds nor the circle, the compound is not aromatic. It is simply cyclohexane and there are two hydrogens on each carbon atom. This is easy to mistake when hurrying, so be careful when you are intepreting any structural formulas which include hexagons. The structure with three double bonds was proposed by Kekule as an attempt to explain how a molecule whose molecular formula was C6H6 could be built out of carbons which make four bonds. The ring and the three double bonds fit the molecular formula, but the structure doesn't explain the chemical behavior of benzene at all well. Each of the double bonds would be expected to show the characteristic behavior of an alkene and undergo addition reactions, but this is not how benzene reacts. In particular, we would expect a carbon-carbon double bond to react quickly with bromine to make a dibromo compound. This is what alkenes do very readily, and in fact it is a useful test for alkenes in the laboratory.

Benzene does not react with bromine unless a very bright light or a strong catalyst is used, and then the reaction is not an addition reaction. We conclude that there is something quite unusual about the double bonds in benzene. Kekule (thinking about this problem before bonds were understood as pairs of electrons) suggested that there are two forms of benzene which differ in the locations of the double bonds. His idea was that these were in rapid equilibrium, so rapid that there was never a fixed location for the double bond. One could say that an approaching bromine molecule could not "find" a double bond to react with. There were several other structures proposed for benzene, but a much more satisfactory approach became possible when we began to understand that covalent bonds consist of pairs of electrons shared between atoms. The difference between the two structures Kekule envisioned (called Kekule structures) is only the difference between the locations of three pairs of electrons. This is exactly the type of situation where resonance must be involved. The hybrid or "average" of the two Kekule structures has one sigma bond and one-half of a pi bond between each two carbon atoms. Thus each carbon is joined to each of its neighbors by a one-and-half bond.

Each bond in the benzene ring has the same number of electrons and is the same length. This picture is in complete accord with experiments which show that all carbon-carbon bonds in benzene are the same length, with no hint of shorter (double) or longer (single) bonds. It also helps explain why benzene does not undergo addition reactions: there are no simple pi bonds. Recall that resonance has another important feature: when resonance is involved, the real structure is more stable than we would expect from any of the structures we write using the one line = two electrons symbolism. This extra lowering of energy, which for benzene is about one-third as much as making a typical covalent bond, is quite important in the reactions of benzene and other aromatic compounds. As we will see, reactions of the benzene ring almost always result in products which in which the benzene ring persists -- an outcome of its stability. When resonance theory was first applied to understanding the structure of benzene, the key feature seemed to be a resonance hybrid of ring structures containing alternating single and double bonds. This immediately led to attempts to make and study compounds like cyclooctatetraene and cyclobutane.

These compounds also have ring structures with alternating single and double bonds. Cyclooctatetraene has been made, but it does not posess the properties of extra stability and resistance to addition reactions which distinquish aromatic compounds. It readily adds bromine, for example. Cyclobutadiene is extremely unstable -- one cyclobutadiene molecule reacts with another cyclobutadiene molecule instantaneously even at very low temperatures -- so it certainly does not act like an aromatic molecule and it has been called "antiaromatic" as a result. It seems that there is more to being aromatic than simply a ring with alternating single and double bonds. After considerable development of the underlying theory, the pattern which has emerged is that aromatic characteristics are only expected when there is a ring of pi electrons in which the number of pi electrons is equal to 4n + 2 (where n is an integer, 0, 1, 2, etc.). (This is known as the Huckel rule after its discoverer.) We can check this against the compounds we have considered so far: Benzene has 6 pi electrons (two for each pi bond) which is the number we get from 4n + 2 if n = 1.

Cyclooctatetraene has 8 pi electrons, and there is no integer "n" which will make 4n + 2 = 8. Cyclobutadiene has 4 pi electrons and also doesn't fit 4n + 2. There are many other examples which support Huckel's rule. It is important to be sure that the ring of alternating single and double bonds is complete. If there is an sp3 hybridized carbon in the ring, the conditions for aromatic character are not present, and we do not worry about checking for 4n + 2. Here's an example: Another way to see this is to look at the p orbitals which combine to make the pi bonds. If these p orbitals combine to form an uninterrupted ring as is the case in benzene, then we can go ahead to use Huckel's rule to check for the proper number of pi electrons for aromatic character. If the ring of p orbitals is broken by a CH2 (group or another tetrahedral carbon) with no p orbital, then the compound cannot be aromatic and we need not try to apply Huckel's rule.

The p orbitals which make up the unbroken p orbital ring can be associated with other atoms than carbon. Two examples are furan and pyrrole, in which two of the six electrons needed come formally from unshared electron pairs on oxygen. Such an unshared pair can also come from a carbon atom, which will have to have a negative charge. An example of this is the cyclopentadienide ion which can be made by treating cyclopentadiene with a moderately strong base. Cyclopentadienide ion is sufficiently stabilized by its aromatic character that cyclopentadiene (its conjugate acid) has a pKa of 16, close to that of water. Cyclopentadiene is a remarkably strong acid for a hydrocarbon because its conjugate base has the extra stability of an aromatic compound. Extraordinarily stable cations can also be made if their structures are aromatic. Here are two: Notice that here the formally positively charged carbon atoms are sp2 hybridized and have an empty p orbital which completes the cyclic arrangement of p orbitals. Let's finish up today by looking at the general mechanism for the characteristic reactions of aromatic compounds -- electrophilic aromatic substitution.

The most important characteristics of these reactions follow directly from the stability of the aromatic ring. First, these reactions are typically catalyzed by strong electrophilic (Lewis acidic) catalysts like H2SO4, AlCl3, and FeCl3 which are required to overcome the stability of the aromatic ring. Second, these are substitution reactions since addition reactions would interrupt the p orbital ring and destroy the aromatic stability. Even though the outcome of the attack of electrophiles on benzene is substitution rather than addition, the first step is the same as in electrophilic addition to alkenes -- attack of the electrophile on a pi bond and the formation of a new sigma bond between a carbon atom and the electrophile. The carbocation which is formed undergoes loss of the H+ from the carbon which was attacked. The electrons from the C-H bond are returned to the aromatic pi electron ring and aromatic stability is restored. Notice that the intermediate here is a carbocation, but it is not aromatic. The carbon bearing the hydrogen and the electrophile is sp3 hybridized and has no p orbital to contribute to a cyclic p orbital system.

The carbocation intermediate is somewhat resonance stabilized, though, by a resonance arrangement which is very similar to the one we saw in the addition of electrophiles to conjugated dienes. This intermediate is a carbocation, and as we will see next time, its stability is important in determining how fast the reaction goes and (in benzene rings which bear substituents at one of the carbons) where the electrophile attacks. The key thing to recognize now is that the positive charge and the corresponding carbocation characteristics only appear at positions ortho and para relative to the point at which the electrophile attack. (Nomenclature is treated in Sec 6.3 of Atkins & Carey.) This will turn out to be quite important, so verify this for yourself. Benzene Series C6H6 Benzene- C6H6 Benzene is hydrocarbon made up of 6 carbons and 6 hydrogen atoms. If forms a six sided ring. It was once believed to be a structure that had alternating double and single bonds. If this were the case, the bond lengths would be longer for the single bonds and shorter for the double. However this is not the case. All the bonds are the same length.

Which means all the bonds of the ring are the same. You can think of the bonds as 1.5 bonds (not a single but not a double). A bond is made up of 2 electrons. Each carbon carbon bond has 3 electrons shared (1.5 bonds). Therefore all the bonds are the same length. Chemists are very efficient, so here is how we draw benzene. History of Benzene ( from Purdue )- In 1825, Michael Faraday was asked to analyze an oily liquid with a distinct odor that collected in tanks used to store coal gas at high pressures. Faraday found that this compound had the empirical formula: CH. Ten years later, Eilhardt Mitscherlich produced the same material by heating benzoic acid with lime. Mitscherlich named this substance benzin, which became benzene when translated into English. He also determined that the molecular formula of this compound is C6H6. Benzene is obviously an unsaturated hydrocarbon because it has far less hydrogen than the equivalent saturated hydrocarbon: C6H14. But benzene is too stable to be an alkene or alkyne.

Alkenes and alkynes rapidly add Br2 to the C=C or CC bonds, whereas benzene only reacts with bromine in the presence of a catalyst: FeBr3. Furthermore, when benzene reacts with Br2 in the presence of FeBr3, the product of this reaction is a compound in which a bromine atom has been substituted for a hydrogen atom, not added to the compound the way an alkene adds bromine. FeBr3 C6H6 + Br2 C6H5Br + HBr Other compounds were eventually isolated from coal that had similar properties. Their formulas suggested the presence of multiple C=C bonds, but these compounds were not reactive enough to be alkenes. Because they often had a distinct odor, or aroma, they became known as aromatic compounds. The structure of benzene was a recurring problem throughout most of the 19th century. The first step toward solving this problem was taken by Friedrich August Kekule in 1865.

(Kekule's interest in the structure of organic compounds may have resulted from the fact that he first enrolled at the University of Giessen as a student of architecture.) One day, while dozing before a fire, Kekule dreamed of long rows of atoms twisting in a snakelike motion until one of the snakes seized hold of its own tail. This dream led Kekule to propose that benzene consists of a ring of six carbon atoms with alternating C C single bonds and C=C double bonds. Because there are two ways in which these bonds can alternate, Kekule proposed that benzene was a mixture of two compounds in equilibrium. Kekule's structure explained the formula of benzene, but it did not explain why benzene failed to behave like an alkene. The unusual stability of benzene wasn't understood until the development of the theory of resonance. This theory states that molecules for which two or more satisfactory Lewis structures can be drawn are an average, or hybrid, of these structures. Benzene, for example, is a resonance hybrid of the two Kekule structures. The difference between the equilibrium and resonance descriptions of benzene is subtle, but important.

In the equilibrium approach, a pair of arrows is used to describe a reversible reaction, in which the molecule on the left is converted into the one on the right, and vice versa. In the resonance approach, a double-headed arrow is used to suggest that a benzene molecule does not shift back and forth between two different structures; it is a hybrid mixture of these structures. One way to probe the difference between Kekule's idea of an equilibrium between two structures and the resonance theory in which benzene is a hybrid mixture of these structures would be to study the lengths of the carbon-carbon bonds in benzene. If Kekule's idea was correct, we would expect to find a molecule in which the bonds alternate between relatively long C C single bonds (0.154 nm) and significantly shorter C=C double bonds (0.133 nm). When benzene is cooled until it crystallizes, and the structure of the molecule is studied by x-ray diffraction, we find that the six carbon-carbon bonds in this molecule are the same length (0.1395 nm). The crystal structure of benzene is therefore more consistent with the resonance model of bonding in benzene than the original Kekule structures.

The resonance theory does more than explain the structure of benzene it also explains why benzene is less reactive than an alkene. The resonance theory assumes that molecules that are hybrids of two or more Lewis structures are more stable than those that aren't. It is this extra stability that makes benzene and other aromatic derivatives less reactive than normal alkenes. To emphasize the difference between benzene and a simple alkene, many chemists replace the Kekule structures for benzene and its derivatives with an aromatic ring in which the circle in the center of the ring indicates that the electrons in the ring are delocalized; they are free to move around the ring. The significance of the circle in the center of this aromatic ring might best be understood by asking: What is wrong with the Kekule structures for benzene? We start by building a sigma-bond skeleton for the benzene ring in which each of the carbon atoms is sp2 hybridized. This leaves us with one electron in a 2p orbital on each of the six carbon atoms. If we assume that the interaction between the 2p orbitals is localized between a pair of carbon atoms, we get one of the Kekule structures for benzene.

Switching the pairs of atoms that form bonds gives us the other Kekule structure. But if we allow the six electrons in the six 2p orbitals to interact to form a set of molecular orbitals, we can delocalize the electrons so that they are free to move from one carbon atom to another around the ring. It is this delocalization of electrons around the aromatic ring that is conveyed by the circle that is often written inside the ring. It is also the delocalization of electrons that makes benzene less reactive than a simple alkene. bonding in benzene - sp2 hybridisation and delocalisation BONDING IN BENZENE Important! This article builds on knowledge about the bonding in methane , and the bonding in ethene. You will find the current page much easier to understand if you read these other ones first. You may also find it useful to read the article on orbitals if you aren't sure about simple orbital theory. You can also read about the evidence which leads to the structure described in this article. That page includes the Kekulé structure for benzene and the reasons that it isn't very satisfactory.

An orbital model for the benzene structure Building the orbital model Benzene is built from hydrogen atoms (1s1) and carbon atoms (1s22s22px12py1). Each carbon atom has to join to three other atoms (one hydrogen and two carbons) and doesn't have enough unpaired electrons to form the required number of bonds, so it needs to promote one of the 2s2 pair into the empty 2pz orbital. So the first thing that happens is . . . Promotion of an electron There is only a small energy gap between the 2s and 2p orbitals, and an electron is promoted from the 2s to the empty 2p to give 4 unpaired electrons. The extra energy released when these electrons are used for bonding more than compensates for the initial input. The carbon atom is now said to be in an excited state. Hybridisation Because each carbon is only joining to three other atoms, when the carbon atoms hybridise their outer orbitals before forming bonds, they only need to hybridise three of the orbitals rather than all four.

They use the 2s electron and two of the 2p electrons, but leave the other 2p electron unchanged. Important! If you have any doubts about this then you should follow the links at the top of the page. The new orbitals formed are called sp2 hybrids, because they are made by an s orbital and two p orbitals reorganising themselves. The three sp2 hybrid orbitals arrange themselves as far apart as possible - which is at 120° to each other in a plane. The remaining p orbital is at right angles to them. Each carbon atom now looks like the diagram on the right. This is all exactly the same as happens in ethene. The difference in benzene is that each carbon atom is joined to two other similar carbon atoms instead of just one. Each carbon atom uses the sp2 hybrids to form sigma bonds with two other carbons and one hydrogen atom. The next diagram shows the sigma bonds formed, but for the moment leaves the p orbitals alone. Remember: A sigma bond is formed by the end-to-end overlap between atomic orbitals. Only a part of the ring is shown because the diagram gets extremely cluttered if you try to draw any more.

Notice that the p electron on each carbon atom is overlapping with those on both sides of it. This extensive sideways overlap produces a system of pi bonds which are spread out over the whole carbon ring. Because the electrons are no longer held between just two carbon atoms, but are spread over the whole ring, the electrons are said to be delocalised. The six delocalised electrons go into three molecular orbitals - two in each. Remember: A molecular orbital is the region of space which contains a bonding pair of electrons. Warning! Be very careful how you phrase this in exams. You must never talk about the p orbitals on the carbons overlapping sideways to produce a delocalised pi bond. This upsets examiners because a pi bond can only hold 2 electrons - whereas in benzene there are 6 delocalised electrons. Talk instead about a "pi system" - or just about the delocalised electrons. In common with the great majority of descriptions of the bonding in benzene, we are only going to show one of these delocalised molecular orbitals for simplicity. In the diagram, the sigma bonds have been shown as simple lines to make the diagram less confusing.

The two rings above and below the plane of the molecule represent one molecular orbital. The two delocalised electrons can be found anywhere within those rings. The other four delocalised electrons live in two similar (but not identical) molecular orbitals. Relating the orbital model to the properties of benzene Note: To get the best out of this section you ought to read the article on the Kekulé structure for benzene. The shape of benzene Benzene is a planar regular hexagon, with bond angles of 120°. This is easily explained. It is a regular hexagon because all the bonds are identical. The delocalisation of the electrons means that there aren't alternating double and single bonds. It is planar because that is the only way that the p orbitals can overlap sideways to give the delocalised pi system. The energetic stability of benzene This is accounted for by the delocalisation. As a general principle, the more you can spread electrons around - in other words, the more they are delocalised - the more stable the molecule becomes. The extra stability of benzene is often referred to as "delocalisation energy".

The reluctance of benzene to undergo addition reactions With the delocalised electrons in place, benzene is about 150 kJ mol-1 more stable than it would otherwise be. If you added other atoms to a benzene ring you would have to use some of the delocalised electrons to join the new atoms to the ring. That would disrupt the delocalisation and the system would become less stable. Since about 150 kJ per mole of benzene would have to be supplied to break up the delocalisation, this isn't going to be an easy thing to do. The symbol for benzene Although you will still come across the Kekulé structure for benzene, for most purposes we use the structure on the right. The hexagon shows the ring of six carbon atoms, each of which has one hydrogen attached. (You have to know that - counting bonds to find out how many hydrogens to add doesn't work in this particular case.) The circle represents the delocalised electrons. It is essential that you include the circle. If you miss it out, you are drawing cyclohexane and not benzene.

Questions to test your understanding If this is the first set of questions you have done, please read the introductory page before you start. You will need to use the BACK BUTTON on your browser to come back here afterwards. A benzene molecule is made up of six carbon atom... - Brainly.com Katie a few seconds ago Hi there! Have questions about your homework? At Brainly, there are 60 million students who want to help each other learn. Questions are usually answered in less than 10 minutes. Try it for yourself by posting a question! :D Katie a few seconds ago Hi there! Have questions about your homework? At Brainly, there are 60 million students who want to help each other learn. Questions are usually answered in less than 10 minutes. Try it for yourself by posting a question! :D What makes benzene molecules planar? - Quora Quora Written Jun 30, 2016 sp2 hybridization of all carbon atoms i.e six carbon atoms in benzene makes it to be planar hexagonal shape. But remember that the ring which we are putting in the benzene is not actually in the centre of the hexagonal.

there will be one electron ring above and below the plane of the hexagon. please see this image. Written Jul 21 According to Valence Shell Electron Repulsion Theory, areas of electron density will repel. A benzene molecule has two areas of electron density, one above and one below the carbon atoms. These form rings due to sp2 hybridisation. As there are no electrons within the molecule so to speak, the shapes of the molecule is planar. 41 Views It's all about the hybridization of the ring. Since all the carbons are sp2 hybridized, it results in a planar structure. 33 Views Written Sep 7 I know, right now, five businesses in my neighborhood that are owned and operated completely by illegal immigrants. Specifically, it’s an auto parts store and four local mechanics that buy from that store. I know this because I used to work at that store. The owner was a friend of mine, and asked me to come on board and help him because business was booming and he trusted me with money. I knew h... Written Sep 30 I came to babysit, and they had 2 biological children and one in foster care. 9 years old.

The foster child just came from a first visit to Dad, in prison. Kid was upset as hell, and couldn’t talk or eat or probably see straight. Foster dad, as I entered to babysit, told me to sit down at the dinner table while foster kid finished his meal. As he didn’t want to eat, I found out fast that that wa... Chemical Reactivity Graphite A comparison of the coronene and corannulene models discloses an interesting difference in their shapes. Coronene is absolutely flat and, aside from the peripheral hydrogens, resembles a layer of graphite. Its very high melting point reflects this resemblance. Corannulene, on the other hand, is slightly curved, resulting in a bowl-like shape. If we extend the structure of corannulene by adding similar cycles of five benzene rings, the curvature of the resulting molecule should increase, and eventually close into a sphere of carbon atoms. The archetypical compound of this kind (C60) has been named buckminsterfullerene because of its resemblance to the geodesic structures created by Buckminster Fuller. It is a member of a family of similar carbon structures that are called fullerenes.

These materials represent a third class of carbon allotropes. Alternating views of the C60 fullerene structure are shown on the right, together with a soccer ball-like representation of the 12 five and 20 six-membered rings composing its surface. Precise measurement by Atomic Force Microscopy (AFM) has shown that the C-C bond lengths of the six-membered rings are not all equal, and depend on whether the ring is fused to a five or six-membered beighbor. By clicking on this graphic, a model of C60 will be displayed. Although C60 is composed of fused benzene rings its chemical reactivity resembles that of the cycloalkenes more than benzene. Indeed, exposure to light and oxygen slowly degrade fullerenes to cage opened products. Most of the reactions thus far reported for C60 involve addition to, rather than substitution of, the core structure. These reactions include hydrogenation, bromination and hydroxylation. Strain introduced by the curvature of the surface may be responsible for the enhanced reactivity of C60. .

Larger fullerenes, such as C70, C76, C82 & C84have ellipsoidal or distorted spherical structures, and fullerene-like assemblies up to C240 have been detected. A model of the C70 fullerene may be examined by clicking here . A fascinating aspect of these structures is that the space within the carbon cage may hold atoms, ions or small molecules. Such species are called endohedral fullerenes. The cavity of C60 is relatively small, but encapsulated helium, lithium and atomic nitrogen compounds have been observed. Larger fullerenes are found to encapsulate lanthanide metal atoms. Interest in the fullerenes has led to the discovery of a related group of carbon structures referred to as nanotubes. As shown in the following illustration, nanotubes may be viewed as rolled up segments of graphite. The chief structural components are six-membered rings, but changes in tube diameter, branching into side tubes and the capping of tube ends is accomplished by fusion with five and seven-membered rings. Many interesting applications of these unusual structures have been proposed.

A model of a nanotube will be displayed by clicking on the diagram Other Aromatic Systems Many unsaturated cyclic compounds have exceptional properties that we now consider characteristic of "aromatic" systems. The following cases are illustrative: Compound Substitution Moderate The first three compounds (cyclic polyenes) have properties associated with alkenes in general. Each reacts readily with bromine to give addition products, as do most alkenes. The thermodynamic change on introducing double bonds into the carbon atom ring is also typical of alkenes (a destabilization of ca. 26 kcal/mol for each double bond). Conjugation offsets this increase in energy by a small amount (4-6 kcal/mol). The remaining four compounds exhibit very different properties, and are considered aromatic. Benzene and pyridine are relatively unreactive with bromine, requiring heat and/or catalysts to force reaction, the result of which is substitution rather than addition. Furan and pyrrole react more rapidly with bromine, but they also give substitution products. This tendency to favor substitution rather than addition suggests that the parent unsaturated ring system has exceptional stability. Thermodynamic measurements support this conclusion.

The enhanced stability, often referred to as aromatic stabilization, ranges (in the above cases) from a low of 16 kcal/mol for furan to 36 kcal/mol for benzene. Factors Required for Aromaticity A planar (or near planar) cycle of sp2 hybridized atoms, the p-orbitals of which are oriented parallel to each other. These overlapping p-orbitals generate an array of π-molecular orbitals. These π-orbitals are occupied by 4n+2 electrons (where n is an integer or zero). This requirement is known as The Hückel Rule. All the aromatic compounds discussed above have 6 π-electrons (n=1). 1,3-Cyclopentadiene and 1,3,5-cycloheptatriene both fail to meet the first requirement, since one carbon atom of each ring is sp3 hybridized and has no p-orbital. Cyclooctatetraene fails both requirements, although it has a ring of sp2 hybridized atoms. This molecule is not planar ( a geometry that would have 135º bond angles ).

Angle strain is relieved by adopting a tub-shaped conformation; consequently, the p-orbitals can only overlap as isolated pairs, not over the entire ring. Furthermore, cyclooctatetraene has 8 π-electrons, a number not consistent with the Hückel Rule. Benzene is the archetypical aromatic compound. It is planar, bond angles=120º, all carbon atoms in the ring are sp2 hybridized, and the pi-orbitals are occupied by 6 electrons. The aromatic heterocycle pyridine is similar to benzene, and is often used as a weak base for scavenging protons. Furan and pyrrole have heterocyclic five-membered rings, in which the heteroatom has at least one pair of non-bonding valence shell electrons. By hybridizing this heteroatom to a sp2 state, a p-orbital occupied by a pair of electrons and oriented parallel to the carbon p-orbitals is created.

The resulting planar ring meets the first requirement for aromaticity, and the π-system is occupied by 6 electrons, 4 from the two double bonds and 2 from the heteroatom, thus satisfying the Hückel Rule. Four illustrative examples of aromatic compounds are shown above. The sp2 hybridized ring atoms are connected by brown bonds, the π-electron pairs and bonds that constitute the aromatic ring are colored blue. Electron pairs that are not part of the aromatic π-electron system are black. The first example is azulene, a blue-colored 10 π-electron aromatic hydrocarbon isomeric with naphthalene. The second and third compounds are heterocycles having aromatic properties. Pyridine has a benzene-like six-membered ring incorporating one nitrogen atom. The non-bonding electron pair on the nitrogen is not part of the aromatic &#960-electron sextet, and may bond to a proton or other electrophile without disrupting the aromatic system. In the case of thiophene, a sulfur analog of furan, one of the sulfur electron pairs (colored blue) participates in the aromatic ring &#960-electron conjugation.

The last compound is imidazole, a heterocycle having two nitrogen atoms. Note that only one of the nitrogen non-bonding electron pairs is used for the aromatic π-electron sextet. The other electron pair (colored black) behaves similarly to the electron pair in pyridine. Annulenes Monocyclic compounds made up of alternating conjugated double bonds are called annulenes. Benzene and 1,3,5,7-cyclooctatetraene are examples of annulenes; they are named [6]annulene and [8]annulene respectively, according to a general nomenclature system in which the number of pi-electrons in an annulene is designated by a number in brackets. Some annulenes are aromatic (e.g. benzene), but many are not due to non-planarity or a failure to satisfy the Hückel Rule. Compounds classified as [10]annulenes (a Hückel Rule system) serve to illustrate these factors.

As shown in the following diagram, 1,3,5,7,9-cyclodecapentaene fails to adopt a planar conformation, either in the all cis-configuration or in its 1,5-trans-isomeric form. The transannular hydrogen crowding that destabilizes the latter may be eliminated by replacing the interior hydrogens with a bond or a short bridge (colored magenta in the diagram). As expected, the resulting 10 π-electron annulene derivatives exhibit aromatic stability and reactivity as well as characteristic ring current anisotropy in the nmr. Naphthalene and azulene are [10]annulene analogs stabilized by a transannular bond. Although the CH2 bridged structure to the right of naphthalene in the diagram is not exactly planar, the conjugated 10 π-electron ring is sufficiently close to planarity to achieve aromatic stabilization. The bridged [14]annulene compound on the far right, also has aromatic properties . A modified [10]annulene, aromatic by nmr criteria, was prepared recently by chemists at California Institute of Technology.

Remarkably, this hydrocarbon is chemically unstable, in contrast to most other aromatic hydrocarbons. To learn more Click Here . Barrelene Formulation of the Hückel rule prompted organic chemists to consider the possible aromaticity of many unusual unsaturated hydrocarbons. One such compound was the 6 &#960-electron bicyclic structure, now known as barrelene. Although the &#960-bonds in barrelene are not coplanar, it was believed that transannular overlap might still lead to aromatic stabilization. A synthesis of barrelene (bicyclo[2.2.2]-2,5,7-octatriene) was accomplished nearly fifty years ago by H. Zimmerman (Wisconsin), using a double Hofmann elimination. As shown in the following diagram, the chemical behavior of this triene confirmed it was not aromatic in the accepted sense of this term. Bromine addition took place rapidly with transannular bond formation, in the same fashion as with norbornadiene (bicyclo[2.2.1]-2,5-heptadiene). Pyrolysis of barrelene gave the expected cycloreversion products benzene and acetylene.

The heat of hydrogenation of barrelene reflects its thermodynamic stability. The value for cyclohexene is -28 kcal/mol, significantly less than one third of the barrelene number. Furthermore, the first double bond of barrelene is reduced with the release of 36.7 kcal/mol heat, indicating destabilization rather than stabilization. The electronic spectrum of barrelene shows a &#960-electron interaction similar to that in related homoconjugated dienes. (λmax ≅220-230 nm). An explanation for the lack of aromatic behavior in the case of barrelene may be found by comparing the orbital symmetry of the six component p-orbitals with those of benzene. Benzene is an annulene in which all six p-orbitals may be oriented with congruent overlapping phases . The cylindrical array of p-orbitals in barrelene cannot be so arranged, as shown in the diagram on the right. There will always be one region (a nodal plane) in which the transannular overlap is incongruent. By clicking on this diagram, a Jmol model of barrelene will be displayed in a separate window. This model may be moved about for viewing.

The p-orbitals of the double bonds may also be displayed. Aromatic Ions Carbanions and carbocations may also show aromatic stabilization. Some examples are: The three-membered ring cation has 2 π-electrons and is surprisingly stable, considering its ring strain. Cyclopentadiene is as acidic as ethanol, reflecting the stability of its 6 π-electron conjugate base. Salts of cycloheptatrienyl cation (tropylium ion) are stable in water solution, again reflecting the stability of this 6 π-electron cation. Antiaromaticity Conjugated ring systems having 4n π-electrons (e.g. 4, 8, 12 etc. electrons) not only fail to show any aromatic properties, but appear to be less stable and more reactive than expected. As noted above, 1,3,5,7-cyclooctatetraene is non-planar and adopts a tub-shaped conformation. The compound is readily prepared, and undergoes addition reactions typical of alkenes. Catalytic hydrogenation of this tetraene produces cyclooctane.

Planar bridged annulenes having 4n π-electrons have proven to be relatively unstable. Examples of 8 and 12-π-electron systems are shown below, together with a similar 10 π-electron aromatic compound. The simple C8H6 hydrocarbon pentalene does not exist as a stable compound, and its hexaphenyl derivative is air sensitive. The 12-π-electron analog heptalene has been prepared, but is also extremely reactive (more so than cyclooctatetraene). On the other hand, azulene is a stable 10-π-electron hydrocarbon that incorporates structural features of both pentalene and heptalene. Azulene is a stable blue crystalline solid that undergoes a number of typical aromatic substitution reactions. The unexpected instability of 4n π-electron annulenes has been termed "antiaromaticity". Other examples may be cited. Thus, all attempts to isolate 1,3-cyclobutadiene have yielded its dimer, or products from reactions with other compounds introduced into the reaction system.

Similarly, cyclopentadienyl cation (4 π-electrons) and cycloheptatrienyl anion (8 π-electrons) show very high reactivity when forced to form. Cyclooctatetraene is a fascinating compound. To see more of its chemistry Click Here .

Question: How many carbon atoms make up a benzene molecule?
Answer: Six